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Concomitant cathodic electro-precipitation during an electrolytic process
Calcareous scaling deposition that takes place naturally due to Ca2+ and Mg2+ content of natural water (water hardness) is a permanent problem. Technical and economic issues pertaining calcareous scaling in industrial plants and domestic equipment have triggered researches on this topic in early ages, including via electrochemical methods which has also been called accelerated scaling process [46-50]. This method implies the reduction of dissolved O2 (Eq. (II.9)) and water (Eq. (II.10)) to favor the mineral deposition on cathode of an electrochemical cell [51, 52].
𝑂2+2𝐻2𝑂+4𝑒−→4𝑂𝐻− (II.9).
2𝐻2𝑂+2𝑒−→𝐻2+2𝑂𝐻− (II.10).
Ever since, the cathodic electrochemical precipitation has been widely documented in the literature and it concerns three major domains of application namely cathodic protection (Table II.4), phosphorus recovery (Table II.5) and water softening (Table II.6). On one hand, the deposition on cathode is deliberate to achieve the intended goal. On the other hand, such as in the case with EAOPs, the cathodic deposition is a major drawback, which would decrease the process efficiency during a long operation. Mineral scaling could result in economic challenges in all kinds of electrochemical reactors and the main influencing parameters are exposed hereinafter including the submillimetric features that will be adopted in this thesis.
Relation between cathode potential, local alkalization and mineral scaling
Corrosion of a metal occurs when the metal (M) spontaneously releases its ion (MZ+) into the solution that it is in contact with. It is a spontaneous reaction when two metal elements are present, with one having more tendency to release electrons (less noble metal) acting as anode. Corroded metal loses its tensile strength, becomes brittle and has short life span. That is the reason why several authors have investigated the corrosion behavior of soft metal and carbon metal that have widely been used in marine infrastructure since they present more economic value. For instance, the formation of corrosion product by applying a corrosion potential has been investigated by Rakitin et al. [53] and Gabrielli et al. [46] on ferrous metal and also by Ben Amor et al. [54], Bousselmi et al. [55] and Marin-Cruz et al. [56] on carbon steel. Furthermore, by applying a more negative potential that corresponds to the potential of the reduction of dissolved O2 or water (Eqs. (II.9)-(II.10)), the reduction reactions producing OH- are accelerated. This continuous production of OH- induces high local pH increase at the cathode vicinity. In the presence of carbonates, Ca2+ and Mg2+ in solution, mineral deposits are therefore formed on the cathode surface (Eqs. (II.11)-(II.13)) [50, 57-59].
𝑀𝑔2++2𝑂𝐻−↔𝑀𝑔(𝑂𝐻)2 (II.11).
𝑂𝐻−+𝐻𝐶𝑂3−↔𝐻2𝑂+𝐶𝑂32− (II.12).
𝐶𝑎2++𝐶𝑂32−↔𝐶𝑎𝐶𝑂3 (II.13).
These precipitates, also referred as scaling, serve as the basis for cathodic protection application, which has vastly been applied in underwater pipelines, marine transportation and above/under water infrastructures. It offers an excellent barrier towards further degradation of metal in contact with water. It significantly reduces the diffusion of dissolved O2 towards the metal surface hence hindering further corrosion.
It has been shown across literature that the kinetics, quantity and crystallography of calcareous deposit are heavily dependent on the applied potential [57, 59-61]. Figure II.12 plots normalized chronoamperometric curves during the formation of calcareous scaling at different applied potentials. According to Figs. II.12(b)-II.12(c) a potential of -1.2 V vs. saturated calomel electrode (SCE) represented a potential threshold [59, 60]. From -0.9 to -1.1 V/SCE, the applied potentials were on the plateau of the O2 reduction producing OH- (Eq. (II.9)) according to the voltammogram plotted in Fig. II.12(a) [62]. The lower the potential, the higher the rate of OH- production, hence faster scaling (Eq. (II.13)) occurred. As a result, cathode surface was completely covered by the passivated film, thus the residual current at the end of the electrolysis was almost nil. At -1.2 V/SCE and below, the applied potential reached the region of water reduction (Eq. (II.10)) [62]. OH- production occurred with concurrent HER. The evolved H2 on the cathode surface induced partial detachment of calcareous scaling [60, 61]. It also led to a porous morphology of the scaling, which made the cathode surface not totally covered by the deposition. Consequently, higher residual current was noticed at the end of electrolysis (Figs. II.12(b)-II.12(d)).
Gas evolution activity on electrode surface
When the applied potentials on cathode and anode surpass the potential of water reduction and oxidation respectively, HER (Eq. (II.10)) and O2 evolution reaction (OER) ((Eq. (II.3)) occur. In the lower region of applied current, i.e., under the condition where HER and OER are already reached, the majority of gases are in dissolved form [63]. On a gas evolving electrode, it has been reported that there are periodic nucleation, growth and detachment of bubbles from electrode surface [63-65]. These periodic evolving and departing gas bubbles have been documented to enhance mass transfer locally on electrode surface [63, 64, 66-68]. The enhancement takes place firstly via microconvection owing to the evolving and departing gas bubbles from the electrode surface (Fig. II.13). Secondly, the displacement of bubbles over the electrode surface induces a mechanism known as forced convection or macroconvection [63-65, 67, 68]. As the applied current density increases, the production of gases becomes more intense. The supersaturation of dissolved gas in the vicinity of electrodes increases too. The adhering gas bubbles could coalesce between each other to produce a film of gas bubbles, so-called gas curtain on electrode surface [66]. It can result in a significant increase in ohmic drop under certain configurations especially when the electrodes are brought close to each other [69-73]. The coalescence of dispersed gas bubbles might contribute towards gas film formation [74]. In such cases, slug flow (no longer dispersion) of bubbles, whose thicknesses were below 5 mm, was reported for example in a duct, [69]. The appearance of slug flow of gas bubbles was fact-driven and even more would be expected under the microfluidic setup [75].
The characterization and mechanistic understanding on electro-precipitation taking place inside submillimetric up to millimetric cells lack studies taking into account the gas evolution reactions in the precipitates stability. This will therefore be one of the objectives of the PhD thesis.
Electrolytic properties and composition
Simple Ca2+ containing solution saturated by carbonate using CO2 bubbling has been used in early stages to investigate calcareous kinetics and nuclei growth mechanism [46-48, 50, 78]. From there onwards, more complex matrices of solutions have been explored (Table II.4) such as simulated groundwater [54, 79, 80], simulated seawater [57-62, 85-87], simulated industrial cooling water [56, 81-83], natural groundwater [55] and real seawater [88, 90, 91] as well as mineral drinking water [77, 84]. This transition from simple matrix into real natural water matrix is necessary to investigate several possible parameters that could alter the kinetics of mineral scaling. While treating real wastewater of multiple origins, the wastewater is often composed of multi-ions having various concentrations depending on their source. It means that several precipitations could occur once their thermodynamic constant of precipitation have been reached [92]. Consequently, cathodic precipitation is done deliberately to retain specific species from the bulk and the retained product could be recovered if it has an added value [93]. A clear example is the phosphorus nutrient recovery. Phosphorus is a nonrenewable resource and, in the meanwhile, the emission of phosphorus into the waste streams represents a rich unrecovered phosphorus source [94]. That is the motivation behind the phosphorus recovery via electrochemical precipitation method which has been attempted as electrochemical phosphorus recovery [95]. This technique is currently under optimization to replace chemical dosing method (e.g. with NaOH or Mg(OH)2 to precipitate PO43- under the form of struvite (MgNH4HnPO43-n) in the waste streams rich with P and N nutrients) [93, 96]. A literature review about the recovery via electro-precipitation routes is provided in Table II.6.
Presence of multi-ions in solution
Among the ionic species which are jointly present in the natural water body, Mg2+ and SO42- are the two most reported species to inhibit the growth of scaling [59-62, 86]. Their inhibiting effect can be clearly seen from Figs. II.15-II.16 in which the chronoamperograms and scanning electron micrographs are respectively shown at varying concentrations of Mg2+ and SO42-. Both chronoamperograms indicated an increase in residual currents as a result of lesser calcareous scaling on electrode surface. They were supported by scanning electron microscopy (SEM) images showing significantly bare electrode surface at the highest Mg2+ and SO42- concentrations. According to literature, Mg2+ hinders the growth of CaCO3 by its incorporation into the CaCO3 crystal lattice [79, 97, 98]. The growth of the most stable CaCO3 crystal (calcite, cubic form as shown in Fig. II.15(b-I)) was then slowed down. It promoted aragonite crystallography (needle-like form as in Figs. II.15(b-II) to II.15(b-IV)) instead, as a result of incorporation of Mg2+ into CaCO3 deposit. As for the SO42- ions, they could impede the CaCO3 growth, either by lowering the cathode local pH [86, 87], adsorbing into the CaCO3 crystal lattice [99], or by increasing the CaCO3 solubility [62, 100].
Presence of scaling/corrosion inhibitors
In the objective to reduce scaling tendency, the addition of inhibitors has been tested. Various studies have been devoted to the kinetic growth of CaCO3 in the presence of inhibitors (Table II.4). The inhibitors could be either organic or inorganic compounds. As the name suggests, the growth of calcareous scaling was impeded in their presence. It can be affirmed by the plots of chronoamperometric and EIS curves in Figs. II.20(a) [84] and II.20(b) [56]. In the former, the curve did not decrease in the presence of optimal scale inhibitor (Figs. II.20(a)). In the latter, high charge transfer resistance (RCT) was measured (diameter of the demi-sphere) in the presence of inhibitors, which described significant corrosion activity of carbon steel (Fig. II.20(b)). Moreover, in the absence of scaling inhibitor, the CaCO3 took typical cubic shape calcite crystallography (Fig. II.20(c) [84]). In the presence of inhibitors, deformed and irregular CaCO3 morphologies have been documented (Figs. II.20(d)-II.20(g) [83, 84]). It has been proposed that the inhibitors inhibited the growth of calcareous scaling by adsorbing or co-precipitating with the calcareous layer [56, 83, 84].
Electrode materials
Studies using different electrode materials as substrate for mineral scaling have been reported in literature and some of them have been listed in Table II.4. The electrode materials have diverse electrochemical responses under current/potential polarization. For example, as indicated in Figs. II.23(a)-II.23(b), the voltammograms between electrode materials were distinct [59, 77]. From Fig. II.23(a), faster O2 reduction into OH- took place in respective order of copper, soft steel and stainless steel and the overpotential for HER were almost identical across all three [77]. In Fig. II.23(b), O2 reduction was firstly observed on gold but the potential window before H2 evolution was larger on gold than on steel [59]. As a result, it was observed in Fig. II.23(c) that the scaling on gold was not yet perturbed by HER gas at similar applied potential compared with the steel [59]. Hence faster, more scaling and less residual current were noticed on gold.
Influence of CO32- towards the mineral electro-precipitation
Electrolysis of electrolyte containing Mg2+ and Ca2+, in presence and absence of CO32-, in the microfluidic flow-by reactor were carried out by applying 0.4 and 4 mA cm-2 and the results are plotted in Fig. IV.5 and Fig. IV.6. From Fig. IV.5, it could firstly be observed that Ca2+ did not precipitate in the electrolyte without carbonate at both current densities. It precipitated only in presence of carbonate. Thus, it has been concluded that Ca(OH)2 was not produced throughout these series of experiment and Ca2+ only electro-precipitated to form CaCO3. This observation can be supported by the value of thermodynamic constant of solubility of Ca(OH)2 that is high (5.00 × 10-6 [22, 23]) – thus highly soluble -, hence the interfacial pH never reached the OHcrit− (i.e. 36.6 mol m-3, pH = 12.6) to precipitate with Ca2+. Secondly, according to Fig. IV.6, Mg2+ only precipitated at 4 mA cm-2 and no Mg2+-based deposition occurred at 0.4 mA cm-2. At low current density, since no decrease of Mg2+ was observed, it can be concluded that Mg(OH)2 was not electro-precipitated. Same remark was given by Deslouis et al. [14] where they did not observe precipitation of Mg(OH)2 when interfacial pH of 9.3 was not yet reached. Our result tells that there was a plateau at 4 mA cm-2 at the very beginning of electrolysis up to 7 min before Mg2+ started to deposit forming Mg(OH)2, as soon as OHcrit− (0.16 mol m-3, interfacial pH = 10.2) was reached (Fig. IV.6). Then, during the first hour of electrolysis at 4 mA cm-2, a slightly higher kinetics of Mg(OH)2 deposition was observed in the electrolyte without carbonate. In the presence of CO32-, there was perhaps a weak competition between Mg(OH)2 and CaCO3 for active site on the surface of cathode [20, 53]. However, it was not accounted for MgCO3, otherwise the kinetics of precipitation with the electrolyte containing CO32- would be faster and not slower than the one without CO32-. To conclude, Mg2+ electro-precipitated independently with regard to carbonate ions.
Moreover, it can be seen in Fig. IV.5 that CaCO3 was already deposited at low current density, unlike Mg(OH)2 even though the thermodynamic constant of solubility of Mg(OH)2 is lower than that of CaCO3. Similar observation was reported by Okstad et al. [54]. The higher initial concentration of Ca2+ – implemented to simulate the composition of reclaimed wastewater from urban WWTPs – can be one of the reasons of this observation, whereby supersaturation degree was already high. On top of that, experimental results showed that 7.2% more CaCO3 deposition occurred at 0.4 mA cm-2 in comparison to higher applied current density of 4 mA cm-2. To our understanding, once CaCO3 electro-precipitation started to occur, it took place regardless of applied current density as long as the current density was equal or higher than the starting point of the nucleation of CaCO3 crystal. In case of low and high applied current densities investigated in this work, CaCO3 already formed at the low japp (0.4 mA cm-2). It would behave similarly at higher applied current density except that within the latter, oxidation and reduction of solvent on the surface of electrodes were occurring at higher rates leading to higher amount of electro-generated gas. Given the fact that at 0.4 mA cm-2 it was already exceeding the limiting current density (0.25 mA cm-2, see Section IV.4.3), gas production was therefore occurring at both applied conditions. It means that the influence of gas evolution was taking place concomitantly with the electro-precipitation on cathode surface [55], but in an antagonist way according to the level of gas production. At 0.4 mA cm-2, the lower evolution of gas on the surface of electrode could enhance the mass transfer of species towards the electrodes [56-59]. This enhancement would be attributed to firstly micro-convection, even at small value of current density [56], due to evolving and departing gas bubbles from the electrode surface. Secondly, forced convection or macro-convection owing to movement and displacement of bubbles over the electrode surface could also occur [57-59]. Therefore, transport of Ca2+ and Mg2+ towards cathode surface for precipitating reaction were enhanced at the low japp condition. This could explain why slightly more CaCO3 was obtained at lower japp. Contrastingly, higher intensity of gas evolution is expected over the surface of cathode at high japp (4 mA cm-2). Therefore, bubble coverage at the electrode surface should have increased as described by Vogt et al. [24]. Therefore, the more intense gas evolution at higher japp could have disturbed the layer of CaCO3 nuclei and in-parallel promoted the detachment of CaCO3 deposit [25, 26, 44]. Consequently, less electro-precipitation of CaCO3 was observed at higher japp.
Furthermore, in presence of Ca2+ (150 mg L-1) and carbonates (TIC of 60 mg-C L-1), SI defined in Eq. (IV.7) equaled to 0.89 in the initial electrolyte. It means that Ca2+ and CO32- concentrations in the bulk were already above their equilibrium state in water (where SI = 0) initially [32]. From the thermodynamic point of view, CaCO3 precipitation in the bulk and scaling on metal substrate could occur. However, the induction time for homogenous and heterogeneous CaCO3 nucleation varies in function of physicochemical component of bulk liquid as well as metal substrate [13, 33, 60]. According to Gabrielli et al., thermodynamics requires a 40-fold supersaturation before the precipitation in the bulk can be observed, i.e., equivalent to SI value around 1.6 [2, 60, 61]. It has also been demonstrated that under low supersaturation conditions, very long induction time (12 h) was required before observing scaling on metal substrate [62]. This might be one of the reasons sluggish kinetics of electro-precipitation of CaCO3 was observed at the very beginning of electrolysis in Fig. IV.5. With respect to the evolution of experimental TIC concentrations (Fig. IV.7(a-d)) as well as to the modeled interfacial CO32- concentrations illustrated in Fig. IV.8, this slow initial kinetics were also noticed. From Fig. IV.8, it can be noticed that on cathode surface, CO32- departed from its initial bulk concentration and rapidly increased owing to electrogenerated OH- once the electrolysis started. With increasing quantity of interfacial CO32-, more CO32- was readily available to react with Ca2+ to form CaCO3 scaling. The supersaturation degree quickly increased, thus promoting the deposition of CaCO3. The concentration of interfacial CO32- kept increasing as long as the rate of production of OH- was superior to the rate of production of CaCO3 (or CO32- consumption at the interface). During this period, more CO32- was shifted from HCO3- form since the OH- concentration also increased, the phenomenon that contributed towards local alkalization. A peak was attained when the rate of OH- production equaled that of CO32- consumed. Beyond this peak, the rate of electro-precipitation of CaCO3 declined due to decreasing concentration of both Ca2+ and CO32- in the electrolyte at later stage of electrolysis as can be observed in Fig. IV.5. When comparing the evolution of interfacial CO32- concentration modeled at 4 mA cm-2 (Fig. IV.8(a)) against 0.4 mA cm-2 (Fig. IV.8(b)), three remarks can be made; firstly, longer time was needed to reach the peak at 𝑟OH− = 𝑟CaCO3. Secondly, smaller concentration of interfacial CO32- was produced at higher current density and thirdly, CaCO3 electro-precipitation occurred at relatively slower rate after the peak of equivalence. The involvement of gas evolution might explain the described behaviors. As above-mentioned, intensified evolving H2 gas on cathode surface at higher japp decreased significantly the effective surface area of cathode even though the OH- production was enhanced in the meantime. Consequently, longer buffer time was observed in Fig. IV.5(a) (inset curve) at the beginning of electrolysis in comparison to Fig. IV.5(b). Lesser amount of interfacial CO32- was also expected at the proximity of cathode surface due to the intervention of evolving gas. This disturbance on cathode took place throughout the electrolysis and consequently, plausible detachment of CaCO3 scaling occurred along the way. In overall, slower kinetics of CaCO3 deposition was observed under high current density, affirmed by the trend depicted in Fig. IV.5.
Table of contents :
I. Introduction
II. State of the art
II.1. Electrochemical microreactors for wastewater treatment application
II.1.1. Influence of applied current density
II.1.2. Influence of flow rate
II.1.3. Influence of interelectrode distance
II.1.4. Influence of electrode material
II.1.5. Influence of initial concentration of pollutant
II.1.6. Design and modularity of microfluidic electrochemical cell
II.1.6.1. Mass transfer in submillimetric electrochemical reactors
II.1.6.2. Microfluidic electrochemical reactor in wastewater treatment application
II.2. Concomitant cathodic electro-precipitation during an electrolytic process
II.2.1. Applied cathode potential/current density
II.2.1.1. Relation between cathode potential, local alkalization and mineral scaling
II.2.1.2. Gas evolution activity on electrode surface
II.2.1.3. Current distribution on gas evolving electrode
II.2.2. Electrolytic properties and composition
II.2.2.1. Presence of multi-ions in solution
II.2.2.2. Presence of organics
II.2.2.3. Presence of scaling/corrosion inhibitors
II.2.2.4. Temperature and pH
II.2.3. Electrode materials
II.2.4. Electrochemical reactor design, scale up and maintenance
II.3. Pharmaceuticals removal by EAOPs
II.4. Concluding remarks
III. Material and methods
III.1. Chemical reagents
III.1.1. Preparation of effluents
III.1.2. Analysis
III.1.3. Electrochemical cell characterization
III.2. Preparation of effluent
III.2.1. Simulated effluents
III.2.2. Reclaimed wastewater effluent (RW)
III.2.3. Electrolyte for electrochemical reactor characterizations
III.3. Experimental setup
III.3.1. Electrochemical system
III.3.2. Pretreatment of electrodes
III.3.3. Post treatment and recovery process
III.4. Analytical methods
III.4.1. Inductively coupled plasma optical emission spectrometry (ICP-OES)
III.4.2. TOC/TN analysis
III.4.3. Chromatography
III.4.3.1. Ionic chromatography
III.4.3.2. High performance liquid chromatography – photo diode array detection (HPLC-PDA)
III.4.4. Spectrophotometry
III.4.5. Scanning electronic microscopy (SEM) coupled to energy dispersive X-ray (EDX) analysis
III.4.6. Electrochemical methods
III.4.6.1. Chronoamperometry and chronopotentiometry
III.4.6.2. Cyclic voltammetry (CV) and linear scan voltammetry (LSV)
III.4.6.3. Electrochemical impedance spectroscopy (EIS)
III.5. Modeling software
III.5.1. Kramers-Kronig test
III.5.2. Zsimpwin®
III.5.3. Aquasim©
III.6. Fitting evaluation criteria between experimental and modeling data
IV. Mineral cathodic electro-precipitation and its kinetic modeling in thin-film microfluidic reactor during advanced electro-oxidation process
IV.1. Introduction
IV.2. Experimental section
IV.3. Modeling
IV.3.1. Electrolyte containing Mg2+ and Ca2+
IV.3.2. Electrolyte containing Ca2+ and CO32-
IV.3.3. Electrolyte containing Mg2+ and CO32-
IV.3.4. Electrolyte containing Mg2+, Ca2+ and CO32-
IV.3.5. Modeling software and fitting evaluation
IV.4. Results and discussion
IV.4.1. Stability of anions in blank solutions using BDD or Pt anode
IV.4.2. Local alkalization on cathode surface: reactions’ selectivity between reduction of dissolved O2 and water
IV.4.3. Influence of electromigration of ionic species
IV.4.4. Influence of matrix of electrolyte
IV.4.4.1. Influence of CO32- towards the mineral electro-precipitation
IV.4.4.2. Influence of Mg2+ towards the mineral electro-precipitation
IV.4.4.3. Influence of Ca2+ towards the mineral electro-precipitation
IV.4.5. Evolution of pH and conductivity
IV.4.6. Mass balance and elements recovery
IV.4.7. Theoretical evolution of Ca2+, Mg2+ and interfacial CO32- during electrolysis in different matrices
IV.5. Conclusions
V. Mass transfer evolution in microfluidic thin film electrochemical reactor: New correlations from millimetric to submillimetric interelectrode distances
V.1. Introduction
V.2. Experimental section
V.2.1. Electrochemical system
V.2.2. Mass transfer characterization
V.3. Results and discussion
V.3.1. Mass transfer behavior over a wide range of interelectrode distances
V.3.2. Mass transfer correlations in microfluidic and millimetric parallel-plate electrochemical reactors
V.4. Conclusions
VI. Role of interelectrode distance and electrogenerated gas bubbles on mineral electro precipitation
VI.1. Introduction
VI.2. Experimental section
VI.3. Modeling
VI.3.1. Kinetics of Mg(OH)2 and CaCO3 electro-precipitation
VI.3.2. Relationship between double layer capacitance, double layer thickness and interelectrode distance
VI.3.3. Modeling and fitting evaluation
VI.4. Results and discussion
VIII
VI.4.1. Kinetics and modeling of mineral electro-precipitation at various interelectrode distances
VI.4.2. Impact of cathode potential on mineral electro-precipitations at different interelectrode distances
VI.4.3. The influence of gas evolution on the formation of mineral electro-precipitation
VI.4.4. Cathode/electrolyte interface impedance study of electro-precipitation at different interelectrode distances
VI.5. Conclusions
VII. Effect of simulated and real wastewaters on the occurrence of electro-precipitation and organic pollutants degradation
VII.1. Introduction
VII.2. Experimental section
VII.3. Results and discussion
VII.3.1. Electro-precipitation in electrolyte consisting of only precipitating elements
VII.3.2. Electro-precipitation in the presence of multi-ions representative of reclaimed wastewater
VII.3.3. Electro-precipitation in the presence of organic matter in simulated wastewater
VII.3.4. Electro-precipitation in simulated versus reclaimed wastewater effluent
VII.3.5. Role of electro-precipitation on the degradation of tylosin as model micropollutant in real wastewater effluent
VII.3.5.1. Evolution of electro-precipitation at different applied current densities
VII.3.5.2. Kinetics and modeling of tylosin degradation as target micropollutant in reclaimed wastewater at different applied current densities
VII.4. Conclusions
VIII. Influence of cathode materials towards the formation of electro-precipitate
VIII.1. Introduction
VIII.2. Experimental section
VIII.3. Results and discussion
VIII.3.1. Influence of the porosity of cathode material on the formation of electro-precipitate .
VIII.3.1.1. Electroactivity of stainless steel, graphite and carbon paper characterized by electrochemical method
VIII.3.1.2. Electro-precipitation on different cathode materials at various applied current densities
VIII.3.1.3. Electro-precipitation on porous cathode during electro-oxidation of reclaimed wastewater
VIII.3.2. Role of H2 evolution overpotential: synergistic effect of interelectrode distance and cathode material to reduce electro-precipitation
VIII.3.2.1. Electro-precipitation on stainless steel and graphite at different interelectrode distances
VIII.3.2.2. Cathode/electrolyte interface study of electro-precipitation on graphite and stainless steel by electrochemical impedance technique
VIII.4. Conclusions
IX. General conclusions
IX.1. General overview
IX.2. Kinetic models of mineral scaling inside microfluidic reactors
IX.3. Mass transfer evolution in submillmetric vs. millimetric configuration
IX.4. Fundamental role of interelectrode distance and energetic performance
IX.5. Technical aspects of electro-oxidation treatment of reclaimed wastewater at different applied current densities
